Buffers bases

• A Bronsted-Lowry acid is a proton donor.
• A Bronsted-Lowry base is a proton acceptor.
• A Lewis acid is an electron pair acceptor.
• A Lewis base is an electron pair donor.
• The ion that causes a solution to become acidic is H+ (hydrogen ion) or, more accurately, H 3 O+ (oxonium ion), as protons react with H 2 O to form it.
• The ion that causes a solution to be alkaline is -OH (hydroxide ion).
• The equation for the ionisation of water is 2H 2 O (l) ⇌ H 3 O + (aq) + - OH (aq) or H 2 O (l) ⇌ H + (aq) + - OH (aq).
• Examples of monobasic acids include HCl and HBr.
• Examples of dibasic acids include H 2 SO 4 and HClO 4.
• Examples of tribasic acids include H 3 PO 4 and H 2 SO 4.
• The equation used to convert concentration of H+ into pH is pH = -log[H+].
• The equation used to convert pH into concentration of H+ is [H+] = 10^-pH.
• A pH scale allows a wide range of H+ concentration to be expressed as simple positive values.
• High pH value means a small [H+].
• If two solutions have a pH difference of 1, the difference in [H+] is a factor of 10.
• The concentration of acid at equilibrium is equal to the concentration of acid after dissociation is assumed when calculating pH of weak acids.
• When an acid is added to a buffer solution, the equilibrium shifts to the left because [H+] increases and the conjugate base reacts with the H+ to remove most of the H+.
• When an alkali is added to a buffer solution, the equilibrium shifts to the right, because [OH-] increases and the small concentration of H+ reacts with OH- to restore the H+ ions.
• The equation used to calculate [H+] of buffer solution is: pH = -log[H+], where [H+] is the concentration of H+ ions in the buffer solution.
• The equation used to calculate [H+] of buffer solution is: pH = -log[H+], where [H+] is the concentration of H+ ions in the buffer solution.
• The buffer system that maintains blood pH at 7.4 is HCO3- and CO2.
• When acid/alkali is added to the buffer system that maintains blood pH at 7.4, H+ + HCO3- ⇌ CO2 + H2O.
• When OH- is added to the buffer system that maintains blood pH at 7.4, H2O + OH- ⇌ HCO3-.
• Titration is the addition of an acid/base of known concentration to a base/acid to determine the concentration.
• An indicator is used to show that neutralization has occurred, as is a pH meter.
• The acid base pairs for the reaction CH 3 COOH + H 2 O ⇌ CH 3 COO - + H 3 O + are Acid 1 Base 2 Base 1 Acid 2.
• A strong acid is a chemical species that dissociates completely in water to produce ions with no remainder.
• Methyl orange is used as an indicator for a strong acid-weak base titration.
• Phenolphthalein is used as an indicator for a strong base-weak acid titration.
• Neither methyl orange or phenolphthalein is suitable as they do not give a sharp change at the end point.
• Methyl orange changes from red in acid to yellow in alkali.
• Phenolphthalein changes from colourless in acid to red in alkali.
• Bromothymol blue changes from yellow in acid to blue in alkali.
• To use a pH metre, remove the pH probe from storage solution and rinse with distilled water.
• Dry the probe and place it into the solution with unknown pH.
• Let the probe stay in the solution until it gives a settled reading.
• Acids dissociate completely.
• Examples of strong acids include Hydrochloric acid, Sulfuric acid, and Nitric acid.
• Concentrated means many mol per dm3, strong refers to amount of dissociation.
• Acids that only partially dissociate are referred to as weak acids.